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D. Atomic and Molecular Structure

D. Atomic and Molecular Structure

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Dissolved Oxygen

Levels in Natural


The dissolved oxygen levels in natural waters are dependent on temperature

and water flow.

• To develop a proper technique for obtaining a natural water sample

• To determine the dissolved oxygen concentration of a natural water sample

• To learn the chemical reactions involved in xing and analyzing a water sample for

dissolved oxygen using the Winkler method


The following techniques are used in the Experimental Procedure:


Streams, rivers, lakes, and oceans play vital roles in our quality of life. They not only

are a source of food supplies with the likes of shrimp and salmon but also provide

recreational opportunities in the forms of boating and swimming. Additionally, the

larger bodies of water such as lakes and oceans affect seasonal weather patterns, producing changes in rainfall and snowfall and generating conditions for hurricanes and


The aesthetic appearance of smaller bodies of water such as rivers and lakes indicates an immediate perception of the quality of the water. Color, surface growth, and

odor are early indicators of the quality of the water and the nature of its marine life. As

the public water supplies of most larger cities rely on the presence of surface water, water

chemists must be keenly aware of the makeup of that water. “How must the water be

treated to provide safe and clean water to the consumers?”

A number of water-quality parameters are of primary interest in analyzing a “natural” water sample: pH, dissolved oxygen, alkalinity, and hardness are but a few. A

quick test, pH, is generally determined with a previously calibrated pH meter; dissolved

oxygen concentrations can be completed with a dissolved oxygen meter (Figure 31.1)

although its availability is less likely than that of a pH meter. Alkalinity and hardness

levels are determined using the titrimetric technique (see Experiments 20 and 21).

The concentration of dissolved oxygen in a water sample is an important indicator of

water quality. Waters with high oxygen concentrations indicate aerobic conditions: clean,

clear, and unpolluted. Low oxygen concentrations indicate anaerobic conditions: high turbidity, foul odors, extensive plant growth on the surface. Dissolved oxygen levels that

drop to less than 5 ppm can stress the existing aquatic life.

The solubilities of oxygen in fresh water (saturated solution) at various temperatures are listed in Table 31.1.


Figure 31.1 Dissolved oxygen

meters can be used for

determining O2(aq) levels in

water samples.

Experiment 31




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Table 31.1 Solubility of Oxygen in Freshwater at Various Temperatures

Temperature (ЊC)







Winkler Method of


ppm O2

Temperature (ЊC)

ppm O2



















The Winkler method of analysis for dissolved oxygen, developed by Lajos Winkler in

1888, is the standard experimental procedure for determining the dissolved oxygen

concentration in water and for the calibration of dissolved oxygen meters.

The Winker test is performed in two parts: (1) the water sample is gathered in the

eld, where the dissolved oxygen is “ xed” with two reagents, and (2) the sample is

titrated for nal analysis in the laboratory within a 48-hour period.

The natural water sample is carefully collected on-site such that no air bubbles

remain trapped in the ask after collection. The oxygen is xed by an immediate

reaction with manganese(II) sulfate in a basic solution:

4 MnSO4(aq) ϩ O2(aq) ϩ 8 NaOH(aq) ϩ 2 H2O(l)

l 4 Mn(OH)3(s) ϩ 4 Na2SO4(aq)


The oxygen is xed as the manganese(III) hydroxide,1 an orange-brown color

precipitate—the more precipitate, the greater is the dissolved oxygen concentration.

While on-site, a basic solution of KI-NaN3 is also added to the sample.2 The

manganese(III) hydroxide oxidizes the iodide ion to the triiodide ion, I3Ϫ, while the manganese(III) reduces to the manganese(II) ion:

2 Mn(OH)3(s) ϩ 3 IϪ(aq) ϩ 6 Hϩ(aq) l I3Ϫ(aq) ϩ 6 H2O(l) ϩ 2 Mn2ϩ(aq)


The resulting solution now has a slight yellow-brown color due to the presence of I3Ϫ


The remainder of the dissolved oxygen analysis is completed in the laboratory (but

within 48 hours). The sample is acidi ed with sulfuric acid to dissolve any precipitate.

A titration of the sample with a standardized sodium thiosulfate solution in the presence

of a starch indicator determines the amount of I3Ϫ generated in the reactions conducted

on-site and provides a direct determination of the dissolved oxygen concentration in the

water sample:

I3Ϫ(aq) ϩ 2 S2O32Ϫ(aq) l 3 IϪ(aq) ϩ S4O62Ϫ(aq)


The starch indicator forms a deep-blue complex with I3Ϫ but is colorless in the presence of IϪ:

I3Ϫ•starch (deep blue) l 3 IϪ ϩ starch (colorless)


From equations 31.1–31.3, 1 mole O2 reacts to produce 4 moles of Mn(OH)3, of which

2 moles of Mn(OH)3 react to produce 1 mole of I3Ϫ. The I3Ϫ, which is the result of the

xing of the dissolved oxygen, reacts with 2 moles of S 2O32Ϫ in the titration.

mol O2 ϭ volume (L) S2O32Ϫ ϫ

mol S2O32Ϫ


1 mol I3Ϫ

L S2O32Ϫ

2 mol S2O32Ϫ

2 mol Mn (OH)2

1 mol O2




1 mol I3

4 mol Mn(OH)3



There is uncertainty among chemists as to the oxidation number of manganese in the precipitate—

MnO(OH)2, the hydrated form of MnO2, often represents the form of the precipitate.


Sodium azide, NaN3, is added to eliminate interference in the dissolved oxygen analysis caused

by the presence of nitrite ion, NO2Ϫ, common in wastewater samples.


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From the data collected and analyzed, the moles of O2 converted to milligrams

divided by the volume of the water sample (in liters) that is titrated results in the dissolved oxygen concentration expressed in mg/L or ppm (parts per million) O2:

mg O2

ϭ ppm O2


L sample

A sodium thiosulfate solution is standardized for the experiment with potassium

iodate, KIO3, a primary standard. In the presence of iodide ion, KIO3 generates a quanti ed concentration of triiodide ion, I 3Ϫ.


IO3Ϫ(aq) ϩ 8 IϪ(aq) ϩ 6 Hϩ(aq) l 3 I3Ϫ(aq) ϩ 3 H2O(l)

This solution is then titrated to the starch endpoint with the prepared sodium thiosulfate solution.


I3Ϫ(aq) ϩ 2 S2O32Ϫ(aq) l 3 IϪ(aq) ϩ S4O62Ϫ(aq)

For the analysis of the dissolved oxygen concentration in a water sample, the standard

Na2S2O3 solution should have a molar concentration of 0.025 M or less.

Standard Solution of

Sodium Thiosulfate 3

Procedure Overview. Three water samples are collected from a source that is

selected either by the student chemist or the instructor. The samples are immediately

“ xed” with the addition of a basic solution of manganese(II) sulfate and a basic solution of KI-NaN3. The samples are stored in the dark on ice and analyzed in the laboratory within ideally 6 hours of sampling. The dissolved oxygen concentrations are

reported in units of parts per million (ppm) O2.

Ask your instructor if a standard solution of Na2S2O3 is available. If so, proceed to

Part B of the Experimental Procedure.



Create and design your own Report Sheet for this part of the experiment.

A. A Standard 0.025 M

Na 2 S 2 O 3 Solution

1. Preparation and standardization of 0.1 M Na2S2O3 solution. Refer to Experiment 29, Parts A and B of the Experimental Procedure for the preparation and

standardization of a 0.1 M Na2S2O3 solution. Prepare only 100 mL of the Na2S2O3

of the solution described in Experiment 29, Part B.1 and standardize the solution

using KIO3 as the primary standard solution (Part B.3–4). Calculate the average

concentration of the Na2S2O3 solution.

2. Preparation of a standard 0.025 M Na2S2O3 solution. Using a pipet and 100-mL

volumetric ask, prepare a 0.025 M Na2S2O3 solution from the standardized 0.1 M


Disposal: Dispose of the test solutions as directed by your instructor.

1. Prepare the ask for sam pling. Thoroughly clean and rinse at least three 250-mL

Erlenmeyer asks and rubber stoppers to t. Allow to air dry.

2. Collect the water sample. Gently lay the ask along the horizontal surface of the

water. See Figure 31.2, page 346. Slowly and gradually turn the ask upright as

the flask fills being careful not to allow any air bubbles to form in the flask.

Fill the ask to over owing.

3. “Fix” the dissolved oxygen. Below the surface of the water sample, pipet ϳ1 mL

of the basic 2.1 M MnSO4 solution into the sample (some over owing will occur).

Similarly pipet ϳ1 mL of the basic KI-NaN3 solution. A precipitate should form

(equation 31.1).

4. Secure the sample.

a. Carefully stopper the sample to ensure that no air bubbles become entrapped

beneath the stopper in the water sample. Again, some over owing will occur.

B. Collection of Water



See Experiment 29 for further explanation and Experimental Procedure.

Experiment 31




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Figure 31.2 Allow a gentle flow of water into the flask.

Slowly turn the flask upright as it fills to overflowing.

b. Invert and roll the ask to thoroughly mix the reagents. Once the precipitate

settles, repeat the mixing process.

c. Label the sample number for each of the asks. Store the sample in the dark

and, preferably, in a cool or cold location or on ice.

5. Temperature. Read and record the temperature of the water at the sample site.

Also, write a brief description of the sample site.

6. Analysis should begin within 6 hours of sampling.

C. Sample Analysis

Read and record the buret to the

correct number of significant figures.

1. Prepare the titrant. Prepare a clean buret. Add 3 to 5 mL of the standard Na2S2O3

solution to the buret, roll the solution to wet the wall of the buret, and dispense

through the buret tip and discard. Use a clean funnel to ll the buret—dispense a

small portion through the buret tip. Read and record the volume of Na2S2O3 solution in the buret (Technique 16A.2), using all certain digits plus one uncertain digit.

Place a white sheet of paper beneath the receiving ask.

2. Prepare sample 1

a. Remove the stopper from the 250-mL Erlenmeyer ask. To the collected water

sample, add ϳ1 mL of conc H2SO4 (Caution!) and stir or swirl to dissolve any

precipitate. The sample can now be handled in open vessels.

b. Transfer a known, measured volume (ϳ200 mL, ‫ע‬0.1 mL) to a receiving ask

(either a beaker or Erlenmeyer ask) for the titrimetric analysis (Part C.3).

3. Titrate water sample 1. Slowly dispense the Na2S2O3 titrant into the water sample. Swirl the ask as titrant is added ( Technique 16C.4). When the color of the

analyte fades to a light yellow-brown, add ϳ1 mL of the starch solution. Continue

slowly adding titrant—when one drop (ideally, half-drop) results in the disappearance of the deep-blue color of the I3Ϫ•starch complex, stop the titration and again

(after ϳ15 seconds) read and record the volume of titrant in the buret.

4. Additional trials. Repeat the analysis for the two remaining samples.

5. Calculations. Calculate the dissolved oxygen concentration for each sample

expressed in ppm O2 (mg O2/L sample ).

Disposal: Dispose of the test solutions as directed by your instructor.

The Next Step


The biological oxygen demand (BOD) of a water sample is a measure of the organic

material in a water sample that is consumable by aerobic bacteria. The O2(aq) concentration is measured when a sample is taken and then again ve days later, that period being

the incubation period for the aerobic bacteria to consume a portion of the O2(aq) to biodegrade the organic material. Research the importance and signi cance of BOD levels in

natural waters and develop an experiment to determine the BOD for a water analysis.

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Experiment 31 Prelaboratory Assignment

Dissolved Oxygen Levels

in Natural Waters

Date __________ Lab Sec. ______ Name ____________________________________________ Desk No. __________

1. For a natural water sample, what range of dissolved oxygen concentrations may you expect? Explain your reasoning.

2. How does the dissolved oxygen concentration in a water sample change (if at all) with

a. ambient temperature changes?

b. atmospheric pressure changes?

c. the volume of the ask collecting the water sample?

d. the amount of organic matter in the water sample?

e. the depth of the body of water (e.g., lake, river, or ocean)?

3. Experimental Procedure, Part B.3. A solution of MnSO4 is added to x the dissolved oxygen in the collected sample.

a. What is the meaning of the expression, “ x the dissolved oxygen,” and why is it so important for the analysis of

dissolved oxygen in a water sample?

b. Only an approximate volume (ϳ1 mL) of MnSO4 is required for xing the dissolved oxygen in the sample. Explain

why an exact volume is not critical.

Experiment 31




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4. A water chemist obtained a 250-mL sample from a nearby lake and xed the oxygen on-site with alkaline solutions of

MnSO4 and KI-NaN3. Returning to the laboratory, a 200-mL sample was analyzed by acidifying the sample with conc

H2SO4 and then titrating with 14.4 mL of 0.0213 M Na2S2O3 solution to the starch endpoint. (This is a calculation

similar to the one for this experiment.)

a. Calculate the number of moles of I3Ϫ that reacted with the Na2S2O3 solution. See equation 31.3.

b. Calculate the number of moles of Mn(OH)3 that were produced from the reduction of the dissolved oxygen. See

equation 31.2.

c. Calculate the number of moles and milligrams of O2 present in the titrated sample. See equations 31.1 and 31.5.

d. What is the dissolved oxygen concentration in the sample, expressed in ppm O2? See equation 31.6.

5. a. Experimental Procedure, Part A. What is the procedure for preparing 250 mL of 0.0210 M Na2S2O3 for this experiment from a 100-mL volume of standard 0.106 M Na2S2O3?

b. For the preparation of the 0.0210 M solution in a 250-mL volumetric ask, only a 25.0-mL calibrated volumetric

pipet is available. Explain how you would prepare the 0.0210 M Na2S2O3 solution using the 25.0-mL pipet. What

would be its exact molar concentration?

6. A 100-mL volume of a primary standard 0.0110 M KIO3 solution is prepared. A 25.0-mL aliquot of this solution is

used to standardize a prepared Na2S2O3 solution. A 15.6-mL volume of the Na2S2O3 solution titrated the KIO3 solution

to the starch endpoint. What is the molar concentration of the Na2S2O3 solution?

IO3Ϫ(aq) ϩ 8 IϪ(aq) ϩ 6 Hϩ(aq) l 3 I3Ϫ(aq) ϩ 3 H2O(l)

I3Ϫ(aq) ϩ 2 S2O32Ϫ(aq) l 3 IϪ(aq) ϩ S4O62Ϫ(aq)


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Experiment 31 Report Sheet

Dissolved Oxygen Levels

in Natural Waters

Date __________ Lab Sec. ______ Name ____________________________________________ Desk No. __________

A. A Standard 0.025 M Na2S2O3 Solution

Prepare a self-designed Report Sheet for this part of the experiment. Review the Report Sheet of Experiment 29 for guidance. Submit this with the completed Report Sheet.

B. Collection of Water Sample

Sampling site: Temperature: ___________ЊC

Characterize/describe the sampling site.

Sample 1

Sample 2

Sample 3

1. Sample volume (mL)




2. Buret reading, initial (mL)




3. Buret reading, nal (mL)




4. Volume Na2S2O3 dispensed (mL)




C. Sample Analysis

5. Molar concentration of Na2S2O3 (mol/L), Part A


6. Moles of Na2S2O3 dispensed (mol)




7. Moles of I3Ϫ reduced by S2O32Ϫ (mol)




8. Moles of O2 (mol)




9. Mass of O2 (mg)







10. Dissolved oxygen, ppm O2 (ppm)

11. Average dissolved oxygen, ppm O2 (ppm)


12. Standard deviation


Appendix B

13. Relative standard deviation (%RSD)


Appendix B

Experiment 31




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Write a short summary based on an interpretation of your analytical data.

Laboratory Questions

Circle the questions that have been assigned.

1. Part B. The water chemist waits until returning to the laboratory to x the water sample for the dissolved oxygen

analysis. Will the reported dissolved oxygen concentration be reported as too high, too low, or remain unchanged?


2. Part B.4. No precipitate forms! Assuming the reagents were properly prepared and dispensed into the sample, what

might be predicted about its dissolved oxygen concentration? Explain.

3. Part B.5. A water chemist measured and recorded the air temperature at 27ЊC when he should have measured the water

temperature, which was only 21ЊC. As a result of this error, will the dissolved oxygen concentration be reported as

being higher or lower than it should be? Explain.

4. Part C.3. The color of the analyte did not fade to form the light yellow-brown color but remained intense even after

the addition of a full buret of the S2O32Ϫ titrant, even though a precipitate formed in Part B.4. What can be stated

about the dissolved oxygen concentration of the sample? Explain.

5. Assuming a dissolved oxygen concentration of 7.0 ppm (mg/L) in a 300-mL water sample,

a. how many moles of Mn(OH)3 will be produced with the addition of the MnSO4 solution?

b. how many moles of I3Ϫ will be produced when the KI-NaN3 solution is added to the above solution?

c. how many moles of S2O32Ϫ will be needed to react with the I3Ϫ that is generated?

d. and also assuming the concentration of the S2O32Ϫ titrant to be 0.025 M, how many milliliters of titrant will be predictably used?

6. A nonscientist brings a water sample to your laboratory and asks you to determine why there was a sh kill in the

nearby lake. Having recently nished this experiment, what might you tell that person about the legitimacy of a test for

dissolved oxygen? What reasoning would you use to maintain the integrity of your laboratory?

7. a. Fish kills are often found near the discharge point of water from cooling waters at electrical generating power

plants. Explain why this occurrence may occur.

b. Fish kills are often found in streams following heavy rainfall in a watershed dominated by farmland or denuded

forestland. Explain why this occurrence may occur.

8. Explain how the dissolved oxygen concentrations may change starting at the headwaters of a river and ending at the

ocean. Account for the changes.

*9. Salt (ocean) water generally has a lower dissolved oxygen concentration than freshwater at a given temperature.

Explain why this is generally observed.


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Galvanic Cells, the

Nernst Equation

Copper metal spontaneously oxidizes to copper(II) ion in a solution

containing silver ion. Silver metal crystals form on the surface of the

copper metal.

• To measure the relative reduction potentials for a number of redox couples

• To develop an understanding of the movement of electrons, anions, and cations in a

galvanic cell

• To study factors affecting cell potentials

• To estimate the concentration of ions in solution using the Nernst equation


The following techniques are used in the Experimental Procedure:


Electrolyic cells are of two types, galvanic and electrolysis, both employing the principle of oxidation–reduction (redox) reactions. In galvanic (or voltaic) cells (this

experiment), redox reactions occur spontaneously as is common with all portable batteries of which we are very familiar. Electric cars, ashlights, watches, and power

tools operate because of a speci c spontaneous redox reaction. Electrolysis cells

(Experiment 33) are driven by nonspontaneous redox reactions, reactions that require

energy to occur. The recharging of batteries, electroplating and re ning of metals, and

generation of various gases all require the use of energy to cause the redox reaction

to proceed.

Experimentally, when copper wire is placed into a silver ion solution (see opening

photo), copper atoms spontaneously lose electrons (copper atoms are oxidized) to the

silver ions (which are reduced). Silver ions migrate to the copper atoms to pick up

electrons and form silver atoms at the copper metal–solution interface; the copper ions

that form then move into the solution away from the interface. The overall reaction that

occurs at the interface is:


Cu(s) ϩ 2 Agϩ(aq) l 2 Ag(s) ϩ Cu2ϩ(aq)

Interface: the boundary between two

phases; in this case, the boundary

that separates the solid metal from the

aqueous solution


This redox reaction can be divided into an oxidation and a reduction half-reaction.

Each half-reaction, called a redox couple, consists of the reduced state and the oxidized state of the substance:

Cu(s) l Cu2ϩ(aq) ϩ 2 eϪ

oxidation half-reaction (redox couple) (32.2)

2 Agϩ(aq) ϩ 2 eϪ l 2 Ag(s)

reduction half-reaction (redox couple) (32.3)

Redox couple: an oxidized and

reduced form of an ion/substance

appearing in a reduction or oxidation

half-reaction, generally associated

with galvanic cells

Experiment 32




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Figure 32.1 Schematic diagram of a galvanic cell

Half-cell: a part of the galvanic cell

that hosts a redox couple

External circuit: the movement of

charge as electrons through a wire

connecting the two half-cells, forming

one-half of the electrical circuit in a

galvanic cell

Salt bridge: paper moistened with a

salt solution, or an inverted tube

containing a salt solution, that

bridges two half-cells to complete the

solution part of an electrical circuit

Internal circuit: the movement of

charge as ions through solution from

one half-cell to the other, forming onehalf of the electrical circuit in a

galvanic cell

Cell Potentials


A galvanic cell is designed to take advantage of this spontaneous transfer of electrons. Instead of electrons being transferred at the interface of the copper metal and the

silver ions in solution, a galvanic cell separates the copper metal from the silver ions to

force the electrons to pass externally through a wire, an external circuit. Figure 32.1 is a

schematic diagram of a galvanic cell setup for these two redox couples.

The two redox couples are placed in separate compartments called half-cells.

Each half-cell consists of an electrode, usually the metal (reduced state) of the redox

couple, and a solution containing the corresponding cation (oxidized state) of the

redox couple. The electrodes of the half-cells are connected by a wire through which

the electrons ow, providing current for the external circuit.

A salt bridge that connects the two half-cells completes the construction of the

galvanic cell (and the circuit). The salt bridge permits limited movement of ions from

one half-cell to the other, the internal circuit, so that when the cell operates, electrical

neutrality is maintained in each half-cell. For example, when copper metal is oxidized

to copper(II) ions in the Cu2ϩ/Cu half-cell, either NO3Ϫ anions must enter or copper(II)

ions must leave the half-cell to maintain neutrality. Similarly, when silver ions are

reduced to form silver metal in its half-cell, either NO3Ϫ anions must leave or cations

must enter its half-cell to maintain neutrality.

The electrode at which reduction occurs is called the cathode; the electrode at

which oxidation occurs is called the anode. Because oxidation releases electrons to the

electrode to provide a current in the external circuit, the anode is designated the negative electrode in a galvanic cell. The reduction process draws electrons from the circuit

and supplies them to the ions in solution; the cathode is the positive electrode. This

sign designation allows us to distinguish the anode from the cathode in a galvanic cell.

Different metals, such as copper and silver, have different tendencies to oxidize; similarly, their ions have different tendencies to undergo reduction. The cell potential of a

galvanic cell is due to the difference in tendencies of the two metals to oxidize (lose

electrons) or of their ions to reduce (gain electrons). Commonly, a measured reduction

potential, the tendency for an ion (or molecule) to gain electrons, is the value used to

identify the relative ease of reduction for a half-reaction.

A potentiometer or multimeter, placed in the external circuit between the two

electrodes, measures the cell potential, Ecell, a value that represents the difference

between the tendencies of the metal ions in their respective half-cells to undergo reduction (i.e., the difference between the reduction potentials of the two redox couples).

Galvanic Cells, the Nernst Equation



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For the copper and silver redox couples, we can represent their reduction potentials as ECu2ϩ,Cu and EAgϩ,Ag, respectively. The cell potential being the difference of the

two reduction potentials is therefore

Ecell ϭ EAgϩ,Ag Ϫ ECu2ϩ,Cu


Experimentally, silver ion has a greater tendency than copper ion does to be in the

reduced (metallic) state; therefore, Agϩ has a greater (more positive) reduction potential. Since the cell potential, Ecell, is measured as a positive value, EAgϩ,Ag is placed

before ECu2ϩ,Cu in equation 32.4.

The measured cell potential corresponds to the standard cell potential when the

concentrations of all ions are 1 mol/L and the temperature of the solutions is 25ЊC.

The standard reduction potential for the Agϩ(1 M)/Ag redox couple, EЊAgϩ,Ag, is

ϩ0.80 V, and the standard reduction potential for the Cu2ϩ(1 M)/Cu redox couple,

EЊCu2ϩ,Cu, is ϩ0.34 V. Theoretically, a potentiometer (or multimeter) would show the

difference between these two potentials, or, at standard conditions,

EЊcell ϭ EЊAgϩ,Ag Ϫ EЊCu2ϩ,Cu ϭ ϩ0.80 V Ϫ (ϩ0.34 V) ϭ ϩ0.46V

Silver jewelry is longer lasting than

copper jewelry; therefore silver has a

higher tendency to be in the reduced

state, a higher reduction potential


Deviation from the theoretical value may be the result of surface activity at the

electrodes or activity of the ions in solution.

In Part A of this experiment, several cells are “built” from a selection of redox couples and

data are collected. From an analysis of the data, the relative reduction potentials for the

redox couples are determined and placed in an order of decreasing reduction potentials.

In Part B, the formations of the complex [Cu(NH3)4]2ϩ and the precipitate CuS are

used to change the concentration of Cu2ϩ(aq) in the Cu2ϩ/Cu redox couple. The

observed changes in the cell potentials are interpreted.

Measure Cell Potentials

The Nernst equation is applicable to redox systems that are not at standard conditions,

most often when the concentrations of the ions in solution are not 1 mol/L. At 25ЊC,

the measured cell potential, Ecell, is related to EЊcell and ionic concentrations by

Measure Nonstandard

Cell Potentials



n log Q

where n represents the moles of electrons exchanged according to the cell reaction. For

the copper–silver cell, n ϭ 2; two electrons are lost per copper atom and two electrons

are gained per two silver ions (see equations 32.1–32.3). For dilute ionic concentrations, the reaction quotient, Q, equals the mass action expression for the cell reaction.

For the copper–silver cell (see equation 32.1):

Nernst equation: Ecell ϭ EЊcell Ϫ



In Part C of this experiment, we study in depth the effect that changes in concentration of an ion have on the potential of the cell. The cell potentials for a number of

zinc–copper redox couples are measured in which the copper ion concentrations are

varied but the zinc ion concentration is maintained constant.

Mass action expression: the product

of the molar concentrations of the

products divided by the product of

the molar concentrations of the

reactants, each concentration raised

to the power of its coefficient in the

balanced cell equation

Zn(s) ϩ Cu2ϩ(aq) l Cu(s) ϩ Zn2ϩ(aq)

The Nernst equation for this reaction is

Ecell ϭ EЊcell Ϫ







Rearrangement of this equation (where EЊcell and [Zn2ϩ] are constants in the experiment) yields an equation for a straight line:

Experiment 32


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